Homework Problems
Chapter 2 Homework Problems: 1, 4, 19, 24, 28 (give the number of
protons, neutrons, and electrons), 30, 38, 41, 44, 46, 54, 64, 66, 67, 72,
76, 82, 84, 96, 101, 104, 114, 116, 117, 124
CHAPTER 2
Atoms, Molecules, and Ions
Early Theories of Matter
The ancient Greeks discussed two possibilities for the essential
property of matter.
Matter is continuous (no “particles” of matter) - Plato, Aristotle,
and a majority of Greek philosophers.
Matter is discrete (composed of particles) - Democritus,
Leucippus, and a small number of Greek philosophers.
Does this process have an end?
yes – particles of matter exist
no – matter is continuous
General Properties of Chemical Systems
As scientists began studying chemical systems they discovered
several general properties of such systems.
1) Conservation of Mass. The total mass in a closed system
remains constant, even if chemical reactions occur.
2) Law of Definite Proportions. All samples of a particular pure
chemical substance contain the same relative amounts of each element
making up the substance.
Examples:
methane
74.9 % C, 25.1 % H
water
88.8 % O, 11.2 % H
copper (II) sulfate
39.8 % Cu, 20.1 % S, 40.1 % O
3) Law of Multiple Proportions. When two elements can
combine to form several different chemical compounds, the ratio of the
amount of the second element combining with a fixed amount of the first
element will be the ratio of small whole numbers.
Example: There are two common compounds of carbon and oxygen
carbon monoxide
1.000 g of C reacts with 1.332 g of O
carbon dioxide
1.000 g of C reacts with 2.664 g of O
g O in carbon dioxide = 2.664 g = 2.000  2
g O in carbon monoxide
1.332 g
1
It is easier to calculate the ratios with the larger number on top,
but that is not required.
Example: The following pure chemical substances can be formed
out of the elements nitrogen and oxygen
nitrogen monoxide
1.000 g of N reacts with 1.142 g of O
nitrogen dioxide
1.000 g of N reacts with 2.285 g of O
nitrous oxide
1.000g of N reacts with 0.5711 g of O
Do these substances demonstrate the law of multiple proportions?
nitrogen monoxide
1.000 g of N reacts with 1.142 g of O
nitrogen dioxide
1.000 g of N reacts with 2.285 g of O
nitrous oxide
1.000g of N reacts with 0.5711 g of O
g O in nitrogen monoxide = 1.142 g = 2.000  2
g O in nitrous oxide
0.5711g
1
g O in nitrogen dioxide = 2.285 g = 4.001  4
g O in nitrous oxide
0.5711 g
1
g O in nitrogen dioxide
g O in nitrogen monoxide
= 2.285 g = 2.001  2
1.142 g
1
So yes, these data are consistent with the law of multiple proportions.
Dalton’s Atomic Theory
A comprehensive theory that accounted for the above observations was proposed by John Dalton, an English chemist, in 1808.
There were three parts to the theory.
1) Elements are composed of particles, called atoms.
a) All atoms of the same element are identical in size, mass,
and chemical properties.
b) Atoms of different elements differ in their size, mass, and
chemical properties.
Dalton’s Atomic Theory (continued)
2) Chemical compounds are composed of atoms of more than one
element.
a) In any particular pure chemical compound the same kinds of
atoms are present in the same relative numbers.
3) Chemical reactions can rearrange atoms, but atoms cannot be
created, destroyed, or converted from atoms of one element to atoms of a
different element.
We now know that some of the hypotheses in Dalton’s atomic
theory are not completely correct; however, the theory represents a good
starting point in understanding the composition of matter.
Consequences of Dalton’s Atomic Theory
Dalton’s theory can be used to explain the observations cited
above.
1) Conservation of mass. Explained by (1) and (3).
2) Law of definite proportion. Explained by (1) and (2).
Example: methane = CH4
Chemical formula - A list of the elements making up a
compound, giving the number of atoms of each element per molecule or
per formula unit of the compound.
3) Law of multiple proportions. Explained by (1) and (2) .
carbon monoxide (CO)
1 C atom: 1 O atom
carbon dioxide (CO2)
1 C atom: 2 O atom
So for a given amount of carbon, carbon dioxide will have twice
as many oxygen atoms (and therefore twice the mass of oxygen) as
carbon monoxide.
Atomic Structure
In Dalton’s atomic theory the smallest particles (atoms) could not
be further broken down. However, a series of experiments, beginning in
the mid-19th century, demonstrated that atoms themselves were
composed of smaller particles.
Radioactivity
In 1895, Antoine Becquerel discovered that some substances
(such as radium and uranium) spontaneously emit “radiation”, a process
called radioactivity.
Three types of radioactivity were found:
alpha () radiation: positively charged particles, now known to
be He2+ nuclei (2 protons + 2 neutrons)
beta () radiation: negatively charged particles, now known to be
electrons
gamma () radiation:
uncharged, now known to
be high energy photons
(particles) of light
Electrons
J. J. Thompson (1897) found that when a high voltage was applied
across two electrodes at low pressure a beam of particles moved from the
negative to the positive electrode. The particles, named electrons, were
negatively charged and the same regardless of the gas between the
electrodes or the metal used in the electrodes. The charge and mass of an
electron were determined experimentally by Millikan (1909).
The “Plum Pudding” Model
Since atoms are electrically neutral, the negative charge of the
electrons in an atom had to be balanced by a positive charge. Thompson
suggested that most of the space within an atom consisted of a positively
charged substance, with electrons embedded within, the “plum pudding”
model.
Rutherford and the Nuclear Atom
To test Thompson’s plum pudding model, Ernest Rutherford
(1909) carried out an experiment where a beam of positively charged
particles (alpha particles) were directed at a thin sheet of gold metal.
The results of this experiment were inconsistent with the plum
pudding model. Rutherford proposed a new model, called the nuclear
model of the atom, that did account for the experimental results.
Subatomic Particles
particle
charge
Coulombs
mass
elementary
kg
amu
proton (p+)
+ 1.60 x 10-19
+1
1.673 x 10-27
1
neutron (n)
0
0
1.675 x 10-27
1
electron (e-)
- 1.60 x 10-19
-1
9.11 x 10-31
0
________
1 amu = 1.6605 x 10-27 kg
mp/me = 1836.
Atomic Structure
nucleus
electron charge cloud
1) The protons and neutrons of the atom are found in a small region in
the center of the atom, called the nucleus. This region contains most of
the mass of the atom, and all of the positive charge.
2) Electrons in the atom form a diffuse cloud of negative charge centered
on the nucleus and occupying most of the volume of the atom.
3) The size of the charge for the proton and electron is the same. The
charge for the proton is positive, and the charge for the electron is
negative. Neutrons have no charge.
4) The type of element for an atom is determined by the number of
protons in the atomic nucleus.
Element (new definition) - An element is a pure chemical
substance composed of atoms, each of which has the same number of
protons in the nucleus.
Hydrogen - one proton per atom
Helium - two protons per atom
Lithium - three protons per atom
.
.
.
.
Similarly, we can now define a compound (new definition) as a
pure chemical substance composed of two or more different kinds of
atoms.
The Periodic Table
Atomic Number and Mass Number
1) The atomic number (Z) is equal to the number of protons in the atom.
2) Since atoms are electrically neutral, the number of electrons in an
atom is also equal to Z, the atomic number.
3) The mass number (A) is equal to the number of protons + neutrons in
the atom.
a) Because protons and neutrons have a mass of approximately 1
(in amu) and electrons have a mass of approximately 0 (in amu) the mass
number is equal to the approximate mass of the atom in amu.
b) Based on the above, the number of neutrons in an atom is
equal to A - Z. So for an atom:
# protons = Z
# electrons = Z
# neutrons = A - Z
Notation For Atoms
We use the following general notation to represent isotopes of
atoms.
mass number
symbol for element
atomic number
Since we can use the symbol for the element and the periodic
table to determine Z, the atomic number, we often omit Z in giving the
symbol for the atom.
Example: 3416S = 34S We can omit the subscript because all
sulfur atoms contain 16 protons.
Isotopes
The atomic number determines the number of protons and
electrons in an atom. This does not place any restrictions on the number
of neutrons in the atom.
It is possible for atoms of the same element to have different
numbers of neutrons. These different types of atoms are called isotopes.
Isotopes of Hydrogen
normal hydrogen
deuterium
tritium
1H
2H
3H
Note that to a very good approximation isotopes of a
particular element are chemically identical to one another.
Example of Notation for Isotopes
As an example of using the above notation, consider the following naturally occurring isotopes of oxygen (Z = 8).
protons
neutrons
electrons
mass number symbol
8
8
8
16
16O
8
9
8
17
17O
8
10
8
18
18O
Example: How many protons, neutrons, and electrons are there in
one atom of 56Fe? What is the approximate mass of one atom of 56Fe in
amu and in kg?
Example: How many protons, neutrons, and electrons are there in
one atom of 56Fe? What is the approximate mass of one atom of 56Fe in
amu and in kg?
# protons = Z = 26
# neutrons = A – Z = 56 – 26 = 30
# electrons = Z = 26 approximate mass (amu) = A = 56
approximate mass (kg) = 56 amu 1.6605 x 10-27 kg
1 amu
= 9.30 x 10-26 kg
Note that the actual mass of one atom of 56Fe is 55.934939 amu.
Atomic Mass Units (amu)
The mass of a single atom of an element, expressed in SI units, is
an extremely small number. For example, the mass of a single atom of
16O is 2.6560 x 10-26 kg. For convenience, we often express values for
atomic mass in terms of atomic mass units (amu).
Atomic mass units are defined as follows
12.00 amu = mass of one atom of 12C (exact)
From this we get 1. amu = 1.6605 x 10-27 kg (approximate)
The mass of any other atom (or particle) is found relative to the
ratio of its mass to the mass of a 12C atom, which can be measured
experimentally.
Mass of particle (amu) =
mass particle • (12.00 amu)
mass 12C atom
Example: A mass spectrometer is a device for determining values
for mass for atoms and molecules. In a particular experiment, the ratio
(mass M/mass 12C) is measured and found to be equal to 7.337. What is
the mass of the molecule M (in amu)?
Example: A mass spectrometer is a device for determining values
for mass for atoms and molecules. In a particular experiment, the ratio
(mass M/mass 12C) is measured and found to be equal to 7.337. What is
the mass of the molecule M (in amu)?
mass M
= 7.337
mass 12C atom
Mass M = 7.337 (mass 12C atom)
= 7.337 (12.00 amu) = 88.04 amu
Note that because of the way we define atomic mass units, the
only isotope whose mass is exactly equal to its mass number is 12C.
isotope
mass (amu)
1H
12C
238U
1.007825
12.000000... (exact)
238.0508
Atomic Mass in the Periodic Table
Because different isotopes of an element have different masses,
the question arises as to which mass should be given in the periodic
table.
For short lived radioactive elements the mass number of the most
stable isotope of the element is listed.
Element
Z
A
technetium (Tc)
43
98
radon (Rn)
86
222
plutonium (Pu)
94
244
Average Atomic Mass
For naturally occurring elements, the value for mass given in the
periodic table is the average atomic mass, based on the natural
abundance of the isotopes that is observed.
In general, we find the average atomic mass as follows:
Mave = f1 M1 + f2 M2 + f3 M3 + … = i=1n fi Mi
where f1, f2,...are the fractions of each isotope observed in nature
M1, M2,…are the corresponding masses for each isotope (in amu)
Note the following
f1 + f2 + f3 + …= 1
fx = % X
100 %
Non-chemical Example
A person has a box of sandwiches. Half of the sandwiches are
6.0 ounces, and half of the sandwiches are 10.0 ounces. What is the
average weight of a sandwich?
Average weight = (0.50)(6.0 oz) + (0.50)(10.0 oz)
= 8.0 ounces
We use the same procedure in finding the average mass of an atom. We
multiply the fraction of each isotope by the mass of that isotope, and then
add the results to find the average mass.
Chemical Example
There are three naturally occurring isotopes of the element
magnesium. Based on the information below, find the average atomic
mass of a magnesium atom.
Isotope
percent
f
M(amu)
24Mg
78.70 %
23.98504
25Mg
10.03%
24.98584
26Mg
11.17%
25.98259
Chemical Example
There are three naturally occurring isotopes of the element
magnesium. Based on the information below, find the atomic mass of a
magnesium atom.
Isotope
percent
f
M(amu)
24Mg
78.70 %
0.7870
23.98504
25Mg
10.03%
0.1003
24.98584
26Mg
11.17%
0.1117
25.98259
So Mave = (0.7870)(23.98504 amu) + (0.1003)(24.98584 amu)
+ (0.1117)(25.98259 amu)
= 24.30 amu, the value given in the periodic table.
Periodic Table
The periodic table is an arrangement of the chemical elements
based on similarities in their physical and chemical properties
The periodic table contains a large amount of useful information
about the chemical elements.
Organization
There are several ways in which the elements in the periodic table
may be classified.
Rows = Periods
Columns = Groups
This is the more important classification.
Elements in the same group usually have similar physical and chemical
properties.
Simplified Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
1A
2A
3A
4A
5A
6A 7A
8A
You are responsible for knowing the names/symbols for elements 1-57,
72-86, and 92.
Major Groups in the Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
Metals, Nonmetals, and Metalloids
Metals:
Usually solid at room temperature (exceptions Cs, Fr, Hg)
Shiny metallic luster
Good conductors of electricity and heat
Malleable (can be hammered into thin sheets)
Ductile (can be drawn into thin wires)
Nonmetals:
Can be solid, liquid, or gas at room temperature
Dull colored (as solids)
Poor conductors of electricity and heat
Not malleable, not ductile
Metalloids (semimetals): Intermediate between metals and nonmetals
Metals, Nonmetals, Metalloids in the Periodic Table
1A
2A
3A
4A
5A
6A 7A
8A
Examples of Elements (as found in nature)
nickel
germanium
(metal)
(metalloid)
sulfur
(nonmetal)
Formation of Ions
Ions are charged particles. Ions can be formed from an atom by
either adding electrons (to form an anion) or removing electrons (to form
a cation). Ions cannot be formed by changing the number of protons in
the atom.
cations
anions
particle
Z
# electrons
particle
Na
11
11
Na+
11
10
Cl-
17
18
Ca
20
20
S
16
16
Ca2+
20
18
S2-
16
18
Cl
Z
17
# electrons
17
Note that the charge of an ion is indicated by a superscript to the right
of the symbol for the ion.
Metals usually form cations, while nonmetals usually form anions.
Just as we can predict the number of protons, neutrons, and
electrons from the symbol for an atom, we can do the same thing for
cations and anions formed from atoms. We do this using the atomic
number (Z) the mass number (A) and the charge of the ion.
Example: How many protons, neutrons and electrons are there
for a 31P3- ion and a 25Mg2+ ion?
Example: How many protons, neutrons and electrons are there
for a 31P3- ion and a 25Mg2+ ion?
31P3-
# protons = Z = 15
# neutrons = A - Z = 31 - 15 = 16
Charge is 3-, so there are 3 more electrons than protons, and so
the number of electrons = 18.
25Mg2+
# protons = Z = 12
# neutrons = A - Z = 25 - 12 = 13
Charge is 2+, so there are two more protons than electrons, and
so the number of electrons is 10.
Ion Charges For Main Group Elements
Main group elements tend to form ions by adding or removing
electrons so that the number of electrons remaining in the ion is equal to
the number of electrons in one atom of the nearest noble gas.
cations (metals)
group 1A (Li, Na, K, Rb, Cs) form 1+ ions
group 2A (Mg, Ca, Sr, Ba) form 2+ ions
group 3A (Al) form 3+ ions
anions (nonmetals)
group 5A (N, P) form 3- ions
group 6A (O, S, Se, Te) form 2- ions
group 7A (F, Cl, Br, I) form 1- ions
Transition Metal Ions
Transition metals form cations. Most transition metals, as well as
a few main group metals like tin (Sn) and lead (Pb), can form ions with
several different charges, but a few transition metals, such as silver,
usually form only one type of cation (Ag+ for silver, Cd2+ for cadmium,
Zn2+ for zinc).
Example:
Iron (Fe)
Fe2+, Fe3+
Copper (Cu)
Cu+, Cu2+
Chromium (Cr)
Cr3+, Cr6+
It is generally not easy to predict which cations a transition metal
will form.
Chemical Formula
The chemical formula for a substance provides information
concerning the composition of the substance.
We can divide substances into two general types.
1) Substances that exist as collections of molecules. In this case
the chemical formula indicates the number of atoms of each elements
present per molecule.
water
(H2O)
phosphorus pentachloride
(PCl5)
nitrous acid
(HNO2)
For organic molecules, the chemical formula is often given in a
way that indicates how the molecule is put together.
ethyl alcohol
dimethyl ether
CH3CH2OH = C2H6O
CH3OCH3 = C2H6O
acetone
CH3COCH3 = C3H6O
Notice that this longer notation makes it possible to distinguish
among different forms (isomers) of organic molecules.
2) Substances that exist as collections of atoms or ions in the
form of a crystal structure, a regular arrangement of the particles making
up the substance.
For these substances, the formula that is given is usually the
empirical formula. An empirical formula gives the relative number of
atoms of each element making up the compound, reduced to the smallest
set of whole number coefficients.
F-
sodium chloride
polonium
NaCl
Po
Ca2+
calcium fluoride
CaF2
Empirical Formula for Molecular Compounds
The empirical formula for a substance gives the relative number
of atoms of each element making up a pure chemical substance, reduced
to the smallest set of integer values. For substances that exist as
molecules, the molecular formula must be an integer multiple of the
empirical formula.
Substance
Chemical formula
Empirical formula
water
H2O
H2O
hydrogen peroxide
H2O2
HO
benzene
C6H6
CH
dichloroethane
C2H4Cl2
CH2Cl
acetic acid
CH3COOH
CH2O
Molecular Compound
A molecular compound is a compound composed of individual
particles called molecules. The bonding between atoms is in such
compounds is due to the sharing of one or more pairs of electrons.
Molecular compounds are usually made up of one or more nonmetallic
elements. Note that these compounds are sometimes called covalent
compounds since the molecules are held together by covalent bonding.
Examples:
HBr
hydrogen bromide
SF6
CS2
carbon disulfide
N2O4 dinitrogen tetroxide
CH3Cl chloromethane
SO2
sulfur hexafluoride
sulfur dioxide
Some elemental substances, such as O2 and N2, exist as individual molecules, but are not compounds, since they are composed of
atoms of a single element.
Ionic Compound
An ionic compound is a compound formed from positive ions
(cations) and negative ions (anions) held together by electrostatic
attraction. For these compounds we usually give the empirical formula,
(sometimes called the formula unit) or smallest electrically neutral
collection of ions making up the compound.
Binary ionic compounds (compounds formed from ions of two
different elements) are usually a combination of a metal cation and a
nonmetal anion.
Examples:
NaCl
sodium chloride (Na+ and Cl-)
Fe2O3
iron (III) oxide (Fe3+ and O2-)
Na2O
sodium oxide (Na+ and O2-)
CuS
copper (II) sulfide (Cu2+ and S2-)
Chemical Formulas For Main Group Ionic
Compounds
Because main group elements form ions with a particular charge
(depending on which group the element is from) we can predict the
chemical formula for a main group ionic compound. We do this by first
finding the charges of the ions formed, and then combining them to get
an overall neutral compound using the smallest set of whole number
coefficients.
Examples: What is the chemical formula for the ionic compound
formed from magnesium and chlorine, from sodium and sulfur, and from
calcium and oxygen?
Examples: What is the chemical formula for the ionic compound
formed from magnesium and chlorine, from sodium and sulfur, and from
calcium and oxygen?
Mg and Cl
Ions are Mg2+ and Cl-, so formula is MgCl2, magnesium chloride.
Na and S
Ions are Na+ and S2-, so formula is Na2S, aluminum sulfide.
Ca and O
Ions are Ca2+ and O2-, so formula is CaO, calcium oxide.
magnesium chloride
sodium sulfide
calcium oxide
Transition Metal Cations
Transition metals can usually form ions with several different
charges (this is also true for a few main group metallic elements like tin
(Sn) and lead (Pb)). While we cannot easily predict which compounds
will form between a transition metal and a main group nonmetal, we can
usually figure out the charge of the transition metal cation if we know the
chemical formula for the compound. This is indicated in the name of the
compound.
Examples:
CuCl
CuCl2
TiO2
MnS2
Examples:
CuCl
Cl- ion, so Cu+ ion
copper (I) chloride
CuCl2
Cl- ion, so Cu2+ ion
copper (II) chloride
TiO2
O2- ion, so Ti4+ ion
titanium (IV) oxide
MnS2
S2- ion, so Mn4+ ion
manganese (IV) sulfide
CuCl
CuCl2
TiO2
Polyatomic Ions
A polyatomic ion is a group of atoms which collectively has a
charge and acts as an ion in an ionic compound.
ion
name
example of ionic compound
NO3-
nitrate ion
NaNO3, Ca(NO3)2, Ni(NO3)2
SO42-
sulfate ion
CuSO4, K2SO4, Al2(SO4)3
Note that when more than one polyatomic ion is present in a
chemical compound the ion is placed in parentheses and the number of
ions per formula unit of compound is given as a subscript outside the
parentheses.
Hydrates
Some ionic compounds can exist in forms where there is a
specific number of water molecules associated with every formula unit
of the ionic compound. Such substances are called hydrates.
Cobalt (II) chloride hexahydrate
Cobalt (II) chloride
Naming Rules For Simple Compounds
Ionic compounds.
a) [Main group metal (1A, 2A, metals and aluminum (3A), Ag+,
Cd2+, Zn2+] + [main group nonmetal]:
name of metal + name of nonmetal + ide
Examples:
K2S potassium sulfide
NaCl sodium chloride
CaI2 calcium iodide
ZnF2 zinc fluoride
Al2O3 aluminum oxide
b) [Transition group metal (or group 4A metal)] + [main group
nonmetal]:
name of metal (charge of metal) + name of nonmetal + ide
Examples:
FeCl2 iron (II) chloride
FeCl3 iron (III) chloride
Cu2S copper (I) sulfide
NiO
PbS2
nickel (II) oxide
lead (IV) sulfide
c) Cation group
NH4+ ammonium ion
Hg22+ mercury (I) ion
Use name of cation group + name of nommetal + ide.
Example:
NH4Br
Hg2Cl2
(NH4)2S
Hg2O
HgO
ammonium bromide
mercury (I) chloride
ammonium sulfide
mercury (I) oxide
mercury (II) oxide
d) Anion group
C2H3O2- is acetate ion
CN- is cyanide ion
OH- is hydroxide ion
N3- is azide ion
CrO42- is chromate ion
MnO4- is permanganate ion
CO32- is carbonate ion
SCN- is thiocyanate ion
C2O42- is oxalate ion
O22- is peroxide ion
Cr2O72- is dichromate ion
ClO3- is chlorate ion
BrO3- is bromate ion
IO3- is iodate ion
NO3- is nitrate ion
SO42- is sulfate ion
PO43- is phosphate ion
+1 oxygen changes the name to per ________ ate
-1 oxygen changes the name to ________ite
-2 oxygen changes the name to hypo ________ ite
Example:
ClO3- is chlorate ion, so
+1 O ClO4- is perchlorate ion
-1 O
ClO2- is chlorite ion
-2 O
ClO- is hypochlorite ion
So we get (use name of cation + name of nonmetal group)
NaClO4
sodium perchlorate
NaClO3
sodium chlorate
NaClO2
sodium chlorite
NaClO
sodium hypochlorite
If more than one of a group is present, we place the group in
parentheses with a number outside indicating how many of the group are
present.
Mg(NO3)2
magnesium nitrate
Hydrogen containing anions
HPO42- is hydrogen phosphate ion
H2PO4- is dihydrogen phosphate ion
HCO3- is hydrogen carbonate (bicarbonate) ion
HSO4- is hydrogen sulfate (bisulfate) ion
Examples:
NaNO3
CuSO4
Zn(ClO3)2
Zn(ClO2)2
KH2PO4
sodium nitrate
copper (II) sulfate
zinc chlorate
zinc chlorite
potassium dihydrogen phosphate
Acids. A substance that produces H+ ions when added to water
(Arrhenius definition).
a) Binary acids (hydrogen + nonmetal)
hydro + nonmetal + ic acid (when in aqueous phase)
Examples:
HBr
H2S
hydrobromic acid
hydrosulfuric acid
b) Ternary acids (hydrogen + oxygen + nonmetal)
hypo ________ite ion becomes hypo ________ ous acid
________ ite ion becomes ________ ous acid
________ ate ion becomes ________ ic acid
per ________ ate ion becomes per ________ ic acid
So -ate is changed to –ic acid, -ite is changed to -ous acid.
Example:
ClOClO2ClO3ClO4-
hypochlorite ion
chlorite ion
chlorate ion
perchlorate ion
HClO hypochlorous acid
HClO2 chlorous acid
HClO3 chloric acid
HClO4 perchloric acid
Binary molecular compounds (compounds are usually two nonmetals)
a) Left or lower element is named first, second element is given
an ide ending, prefix is used to indicate the number of atoms per
molecule (but the prefix mono is never used for the first element)
mono = 1
tri = 3
penta = 5
hepta = 7
di = 2
tetra = 4
hexa = 6
octa = 8
Example:
NO
nitrogen monoxide
NO2 nitrogen dioxide
N2O dinitrogen monoxide (nitrous oxide)
b) A few molecules have common names: H2O is water, NH3 is
ammonia, CH4 is methane.
c) If a binary molecular compound forms an acid when added to
water the naming of the compound depends on whether it is in the gas
phase or aqueous phase.
HCl(g)
HCl(aq)
hydrogen chloride
hydrochloric acid
Organic Molecules
There are systematic rules for naming organic molecules, but you
will not be responsible for naming such compounds.
CH3CH2CH2CH2CH3
CH3CH2OH
CH3CH2OCH2CH3
n-pentane
ethyl alcohol
diethyl ether
Hydrocarbons
Hydrocarbons are molecules that contain only carbon and
hydrogen. One kind of hydrocarbon in an alkane. Alkanes have the
general formula CnH2n+2, where n = 1, 2, 3, …
n
Formula
Name
1
2
3
4
5
•
•
•
CH4
C2H6
C3H8
C4H10
C5H12
•
•
•
methane
ethane
propane
butane
pentane
•
•
•
Functional Groups
Organic molecules can also be classified based on whether or not
they contain a particular group of atoms, called a functional group.
Molecules containing the same functional group often have similar
physical and chemical properties.
End of Chapter 2
“…the ultimate particles of all homogeneous bodies are perfectly
alike in weight, figure, and so forth.”
- John Dalton, A New System of Chemical Philosophy (1808)
“Elements arranged according to the size of their atomic weights
show clear periodic properties.” - D. I. Mendeleev (1869)
“I don’t believe that atoms exist!” - Ernst Mach (1897)
“One of the wonders of this world is that objects so small can
have such consequences: Any visible lump of matter - even the merest
speck - contains more atoms than there are stars in our galaxy.”
- P. W. Atkins